Galvanic Cell Notation Example - University of Missouri ...

1 Galvanic Cell Notation Inactive (inert) electrodes not involved in the electrode half-reaction (inert solid conductors;. Half-cell Notation serve as a contact between the Different phases are separated by vertical lines solution and the external el. circuit). Species in the same phase are separated by Example : Pt electrode in Fe3+/Fe2+ soln. commas Fe3+ + e- Fe2+ (as reduction). 3+ 2+. Notation : Fe , Fe Pt(s). Types of electrodes Active electrodes involved in the electrode Electrodes involving metals and their half-reaction (most metal electrodes) slightly soluble salts Example : Zn2+/Zn metal electrode Example : Ag/AgCl electrode Zn(s) Zn2+ + 2e- (as oxidation) AgCl(s) + e- Ag(s) + Cl- (as reduction). Notation : Zn(s) Zn2+ Notation : Cl- AgCl(s) Ag(s). Electrodes involving gases a gas is bubbled Example : A combination of the Zn(s) Zn2+ and over an inert electrode Fe3+, Fe2+ Pt(s) half- cells leads to: Example : H2 gas over Pt electrode H2(g) 2H+ + 2e- (as oxidation).

2 Notation : Pt(s) H2(g) H+. Cell Notation The anode half-cell is written on the left of the cathode half-cell Zn(s) Zn2+ + 2e- (anode, oxidation). The electrodes appear on the far left (anode) and +. Fe3+ + e- Fe2+ ( 2) (cathode, reduction). far right (cathode) of the Notation Zn(s) + 2Fe3+ Zn2+ + 2Fe2+. Salt bridges are represented by double vertical lines Zn(s) Zn2+ || Fe3+, Fe2+ Pt(s). 1. Example : A combination of the Pt(s) H2(g) H+ Example : Write the cell reaction and the cell and Cl- AgCl(s) Ag(s) half- cells leads to: Notation for a cell consisting of a graphite cathode Note: The immersed in an acidic solution of MnO4- and Mn2+. reactants in the and a graphite anode immersed in a solution of Sn4+. overall reaction are and Sn2+. in different phases Write the half reactions (a list of the most common (no physical half-reactions is given in Appendix D).)

3 Contact) no need Sn2+ Sn4+ + 2e- 5 (oxidation). +. of a salt bridge MnO4 + 8H + 5e Mn + 4H2O(l) 2 (reduction). - + - 2+. H2(g) 2H+ + 2e- (anode, oxidation) 5Sn2+ + 2 MnO4- + 16H+ + 10e- 5Sn4+ + 10e- +. +. AgCl(s) + e- Ag(s) + Cl- ( 2) (cathode, reduction) + 2Mn2+ + 8H2O(l). 2 AgCl(s) + H2(g) 2Ag(s) + 2H+ + 2Cl- The graphite (C) electrodes are inactive Pt(s) H2(g) H+, Cl- AgCl(s) Ag(s) C(s) Sn2+, Sn4+ || H+, MnO4-, Mn2+ C(s). Why Do Galvanic cells Work? Cell Potentials Consider a cell made of two active metal Electromotive force (emf) drives the electrodes, M1 and M2, and their ions. electrons in the el. circuit If the cell circuit is open, the two metals are in equilibrium with their ions emf is the difference between the electrical potentials of the two electrodes (voltage).

4 1) M1 M1+ + e- 2) M2 M2+ + e- The produced electrons accumulate in the metal Cell potential (Ecell) Ecell = emf electrodes and produce electrical potentials Units volts (V) (1 V = 1 J/C since the If M1 has a greater tendency to give out its electrons, electrical work is equal to the applied voltage the 1st equilibrium is shifted further to the right and times the charge moving between the electrodes). the potential of M1 is more negative When the circuit is closed, electrons flow from the Standard cell potential (Eocell) the cell more negative M1 (anode) toward the less negative potential at standard-state conditions (gases . M2 (cathode) 1 atm, solutions 1 M, liquids & solids pure). 2. Ecell is measured with a voltmeter Electrode potentials (E) characterize the If the (+) terminal of the voltmeter is connected individual electrodes (half-reactions).

5 To the (+) electrode (cathode), the voltmeter The cell potential is the difference between the shows a positive reading electrode potentials of the cathode and anode Ecell characterizes the overall cell reaction Ecell = Ecathode Eanode If Ecell > 0, the cell reaction is spontaneous Standard electrode potentials (Eo) . If Ecell < 0, the cell reaction is non-spontaneous electrode potentials at the standard-state If Ecell = 0, the cell reaction is at equilibrium Eocell = Eocathode Eoanode Example : Zn(s) Zn2+(1M)|| Cu2+(1M) Cu(s) Eo values are reported for the half-reaction + V written as reduction (standard reduction potentials) listed in Appendix D. Zn(s) + Zn2+ + Cu(s). Cu2+. Eocell = V > 0 spontaneous reaction Absolute values for E and Eo can't be measured If the unknown electrode is the cathode of the cell A reference electrode (half-cell) is needed Eocell = Eounkn Eoref The potentials of all electrodes are measured relative Eounkn = Eocell + Eoref = Eocell + 0 = Eocell > 0.

6 To the reference electrode If the unknown electrode is the anode of the cell Standard hydrogen electrode used as a Eocell = Eoref Eounkn reference electrode Eoref = 0 V (assumed) Eounkn = Eoref Eocell = 0 Eocell = Eocell < 0. H+(1M) H2(g, 1atm) Pt(s) Example : 2H+(1M) + 2e- H2(g, 1atm) Pt(s) H2(g, 1atm) H+(1M), Cl-(1M) AgCl(s) Ag(s). To find the potential of any electrode, a cell is H+/H2 anode constructed between the unknown electrode and Ag/AgCl cathode the reference electrode Eocell = EoAg/AgCl Eref The cell potential is directly related to the = EoAg/AgCl unknown electrode potential EoAg/AgCl = + V. 3.